Download Bonding Intro 02-09

Ch 2 Structure & Bonding
An amazing thing about the universe - It works in a way that
sometimes when things come together, they stick…
Lewis Structure Reminders
Lots of practice!
Work all the examples you can find!
• Protons and neutrons in a atomic nuclei
• Atoms in molecules
H
H
H
H
“spinspin-pairing”
Hints on Lewis Dot Structures
1. Octet rule is the most useful guideline.
2. Carbon forms 4 bonds.
3. Hydrogen typically forms one bond to other
atoms.
4. When multiple bonds are forming, they are
usually between C, N, O or S.
5. Nonmetals can form single, double, and
triple bonds, but not quadruple bonds.
6. Always account for single bonds and lone
pairs before forming multiple bonds.
7. Look for resonance structures.
“Optimal” resonance structures:
Never violate octet rule for n=2
Satisfy octet rule whenever possible
Can violate octet rule for n > 2
Reduce formal charge if possible
No negative charge on low c atoms
BTW: A Note on Formal Charges
#Valence e-’s - # Actual e-’s
⎡1
⎤
FC = G.N. - ⎢ # BE + # LPE ⎥
⎣2
⎦
“If Step 4 leads to a positive formal charge on an
inner atom beyond the second row, shift
electrons to make double or triple bonds to
minimize formal charge, even if this gives an
inner atom with more than an octet of electrons.”
1
All resonance structures are not
created equal!!!
Chemical Bonds: attractive force holding
two or more atoms together
TWO EXTREME CASES
• Covalent bonding: when atoms share electrons (e.g.
“organic” carboncarbon-based molecules)
• Ionic bonding: when “Atom A” transfers an electron to
“Atom B” (e.g. salts such as NaCl)
NaCl)
• PolarPolar-covalent bonding: everything in between these
two extremes (a good portion of ‘reality”
KEY POINT: Bonding often happens so all bonding
partners acquire a noblenoble-gas electron configuration
Covalent Bonding: Considerations
Covalent Bonding: Considerations
Electrons & nuclei “balance
“balance”” all interactions to
give the molecule stability.
Balance achieved when electrons are
concentrated between the nuclei (electron sharing)
A covalent bond forms IF:
IF:
„
the energy *released* from sharing electrons is
greater than (>)
Each electron in bonding pair has greater space
available than in the unbonded individual atoms, and
each gets to “feel” the positive charge of both nuclei.
„
the energy *spent* in partially removing one or more
electrons from each of contributing atom
2
Atomic size & shape affect bonding!
& symmetry!!!
Unequal Electron Sharing
A pure covalent bond occurs only when
two identical atoms are bonded: N2
Polar Covalent Bond:
Bond: Unequal sharing
between two dissimilar atoms
„
„
Therefore, the electrons are nearer to one of
the atoms, and that atom acquires a partial
negative charge (δ−).
And consequently the other atom has a partial
positive charge (δ+).
Bond is referred to as polar & the molecule can
be called a dipole (having two poles
KEY QUESTION: How do you determine which atom
has the partial negative charge and which atom has
the partial positive charge?
3
Electronegativity is the Answer!
p. 58
All atoms are not created equal!
Defintion:
Defintion: Electronegativity (χ)
The extent to which an element attracts an
electron toward itself
The bigger the EN difference, the more
polar the bond.
Trends similar to electron affinity
Ionic vs. Covalent
Use definition of IE to define “ionic”
∆χ > 1.7 : Ionic
(H is never ionic)
∆χ < 1.7 : Polar covalent
e.g. : C - Cl versus K - Cl
∆χ = 0.5
∆χ = 2.2
What affects bond length?
1. The smaller the principle quantum numbers
of the valence orbitals, the shorter the bond.
2. The higher the bond multiplicity, the shorter
the bond.
3. The higher the effective nuclear charge of
the bonded atoms, the shorter the bond.
4. The larger the electronegativity difference,
the shorter the bond.
4
Bond Energy
Oxidation Numbers
1. Bond strength increases as more
electrons are shared between the atoms
2. Bond strength increases as the
electronegativity difference (∆
(∆χ) between
bonded atoms increases.
3. Bond strength decreases as bonds
become longer.
Valence Bonding (Localized) vs.
Molecular Orbital Theory (Delocalized)
Orbital Overlap (Localized Bonding)
Bonding orbitals are constructed by combining atomic
orbitals from adjacent atoms.
From Quantum Mechanics: orbitals can add or
subtract; therefore constructive or destructive
interference is possible
5
Localized bonding… H2
Orbital Overlap: As two H atoms approach, the overlap
of their 1s atomic orbitals increases. The wave amplitudes
add, generating a new orbital with high electron density
between the nuclei.
Oh shoot… What about methane?
PH3
Phosphine is a colorless, highly toxic gas
with bond angles of 93.6°
93.6°. Describe the
bonding in PH3.
s and p hybridization
What is the electron
configuration of methane?
6
Methane hybridization
General Features of Hybridization
1.
2.
3.
4.
The # of valence orbitals generated by hybridization
equals the # of valence AOs participating in
hybridization.
The steric number of an inner atom uniquely
determines the number and type of hybrid orbitals.
Hybrid orbitals form localized bonds by overlap with
atomic orbitals or with other hybrid orbitals.
There is no need to hybridize orbitals on outer atoms,
because atoms do not have limiting geometries.
The bonds formed by all other outer atoms can be described
using valence p orbitals.
7